Study of Compounds Class 10 ICSE — Chemistry Chapter 10

Written by Greya LakshmiPublished on 3 October 2026

Study of Compounds Class 10 ICSE — the nitric acid chapter in your Class 10 Chemistry book — covers the acid from its discovery to its tests. You will meet aqua fortis, the laboratory preparation, the physical and chemical properties, the uses, and the reactions that identify nitric acid and its salts. Each section below follows the chapter in the book’s own order, with every equation explained.

Download Study of Compounds Nitric Acid Class 10 ICSE Chemistry Chapter 10 PDF

Download the Study of Compounds Nitric Acid Class 10 ICSE Chemistry chapter PDF to revise offline or print for practice. It is the full chapter as printed in the book — 7 pages, 7 sections, 25 questions.

What is inside Chapter 10

7 pages · 7 sections · 25 exercise questions

Section	Title
10.1	INTRODUCTION
10.2	DISCOVERY
10.3	OCCURRENCE
10.4	LABORATORY PREPARATION OF NITRIC ACID
10.5	PROPERTIES OF NITRIC ACID
10.6	USES OF NITRIC ACID
10.7	TESTS FOR NITRIC ACID AND NITRATES

Intext Questions: 11 questions
EXERCISE: 14 questions

10.1 INTRODUCTION

Nitric acid was formerly known as aqua fortis, meaning “strong water.” It earned that name because it reacts with nearly all metals. Silver is the striking case: nitric acid dissolves it, while other acids do not.

The name points to the property that runs through the whole chapter. Nitric acid is not a timid acid. Drop copper turnings into it and the metal is attacked and dissolves — a reaction the chapter returns to later. That aggressive behaviour comes from two features studied in this ICSE chapter on compounds: it is a strong acid, and it is a powerful oxidising agent.

Both features are worked through in section 10.5. The tests in section 10.7 identify the acid and its salts by the brown fumes they release. Nearly all metals are attacked, but not all: gold and platinum resist nitric acid alone, and aqua regia is the answer to that problem.

10.2 DISCOVERY

Three scientists built the modern picture of nitric acid over more than a century. Their work forms a clean timeline, because each step identified one more element in the compound.

Scientist	Year	Finding
Glauber	1658	Obtained nitric acid by distilling nitre (potassium nitrate, $KNO_3$ ) with sulphuric acid
Lavoisier	1776	Proved that nitric acid contains oxygen
Cavendish	1784	Proved that hydrogen and nitrogen are present too

Glauber first made the acid. Lavoisier then proved it contains oxygen. Cavendish completed the analysis by showing hydrogen and nitrogen are also present. Each step removed one unknown, and by 1784 the composition was settled.

Glauber’s method — distilling a nitrate with sulphuric acid — is still the basis of the preparation in section 10.4. The study of compounds chapters in this Class 10 ICSE book cover ammonia, nitric acid, and sulphuric acid in turn.

10.3 OCCURRENCE

Nitric acid occurs in two states. In the free state, it appears in traces in rainwater after lightning. In the combined state, it is found as metallic nitrates: Chile saltpetre ( $NaNO_3$ ), nitre ( $KNO_3$ ), and calcium nitrate $[Ca(NO_3)_2]$ . These three salts matter because the laboratory preparation in section 10.4 starts from them.

The free acid forms through a three-step chain in the atmosphere. Lightning supplies the energy that forces nitrogen and oxygen to combine.

Step 1: nitrogen and oxygen form nitric oxide.

N_2 + O_2 \rightleftharpoons 2NO

Step 2: nitric oxide is oxidised to nitrogen dioxide.

2NO + O_2 \rightleftharpoons 2NO_2

Step 3: nitrogen dioxide dissolves in atmospheric moisture or rain and, with oxygen, forms nitric acid.

4NO_2 + 2H_2O + O_2 \rightarrow 4HNO_3

The rain washes the acid down, and it combines with salts on the soil surface. This conversion of free atmospheric nitrogen into useful nitrogenous compounds is called fixation of atmospheric nitrogen.

10.4 LABORATORY PREPARATION OF NITRIC ACID

Reactants: concentrated sulphuric acid with potassium nitrate (nitre) or sodium nitrate (Chile saltpetre), mixed in equal parts by weight. The mixture is heated gently to $180-200^\circ C$ in a glass retort. Sulphuric acid is non-volatile, so it displaces the volatile nitric acid from the nitrate.

$KNO_3 + H_2SO_4 \xrightarrow{\leq 200^\circ C} KHSO_4 + HNO_3$
$NaNO_3 + H_2SO_4 \xrightarrow{\leq 200^\circ C} NaHSO_4 + HNO_3$

The vapours are condensed by chilling the receiver with running cold water. The collected acid is light yellow, though pure nitric acid is colourless.

The yellow colour has a simple cause. Some of the acid decomposes from the heat, and the reddish-brown nitrogen dioxide produced dissolves in the liquid:

4HNO_3 \rightarrow 2H_2O + 4NO_2 + O_2

The colour is removed in two ways.

Bubble dry air or $CO_2$ through the warm acid. The gas drives out the $NO_2$ , which is then oxidised back to nitric acid.
Add excess water. The $NO_2$ dissolves in the water, reversing the decomposition.

Three precautions protect the preparation. Use an all-glass apparatus, because the vapours attack rubber and cork. Never use concentrated HCl in place of sulphuric acid, since HCl is volatile and its vapours would contaminate the nitric acid.

Keep the temperature at or below $200^\circ C$ . Above it, sodium sulphate forms a hard crust that sticks to the retort, the acid decomposes, the glass can be damaged, and fuel is wasted.

10.5 PROPERTIES OF NITRIC ACID

(A) Physical properties. Pure nitric acid (98%) is a colourless liquid; commercial acid (68%) is yellowish brown. It has a suffocating smell, an acidic taste, and is hygroscopic, so it fumes in air and its bottle must be stoppered. Its density is $1.54\ \text{g/cm}^3$ at 98%, it boils at $86^\circ C$ , melts at $-42^\circ C$ , and it dissolves in water in all proportions. A 68% solution forms a constant boiling mixture at $121^\circ C$ .

It is non-poisonous but corrosive to the skin, causing painful blisters. It reacts with skin protein to form xanthoproteic acid, which leaves a yellow stain.

(B) Stability. Pure nitric acid is unstable to heat or sunlight and decomposes, turning yellow as $NO_2$ forms:

$4HNO_3 \xrightarrow[katex display="true"]\text{Sunlight}$ {\Delta} 4NO_2 + 2H_2O + O_2[/katex]

In sunlight the decomposition happens even at room temperature, so the acid is stored in coloured bottles.

Acidic character. Nitric acid is a strong monobasic acid, ionising almost completely in water:

HNO_3(aq) \rightleftharpoons H^+ + NO_3^-

It turns blue litmus red and methyl orange pink, and leaves phenolphthalein colourless. It neutralises alkalis to salt and water. It reacts with carbonates and bicarbonates to give salt, water, and $CO_2$ . With metallic sulphites and bisulphites it gives the nitrate, water, and $SO_2$ .

Oxidising properties. Nitric acid is a powerful oxidising agent. The power comes from nascent oxygen released as the acid decomposes. Concentrated acid gives $NO_2$ ; dilute acid gives NO:

$2HNO_3\ (\text{conc.}) \rightarrow 2NO_2 + H_2O + [O]$
$2HNO_3\ (\text{dil.}) \rightarrow 2NO + H_2O + 3[O]$

Here is the trap in this study of compounds chapter. Dilute nitric acid does not give hydrogen with metals, as a typical acid would, because the nascent oxygen oxidises the hydrogen to water. Copper with cold dilute acid gives copper nitrate and NO; with concentrated acid it gives $NO_2$ . The one exception is very dilute (about 1%) acid with magnesium and manganese at room temperature, which does give hydrogen:

$Mg + 2HNO_3 \rightarrow Mg(NO_3)_2 + H_2$
$Mn + 2HNO_3 \rightarrow Mn(NO_3)_2 + H_2$

Iron, aluminium, cobalt, and nickel become passive in pure concentrated acid: a thin insoluble oxide layer stops the reaction. Rubbing with sandpaper or treating with a strong reducing agent removes the passivity.

Aqua regia. One part concentrated nitric acid with three parts concentrated hydrochloric acid forms aqua regia, “royal water.” The nitric acid oxidises the HCl to nascent chlorine:

HNO_3 + 3HCl \rightarrow NOCl + 2H_2O + 2[Cl]

The nascent chlorine dissolves noble metals that nitric acid alone cannot:

Au + 3[Cl] \rightarrow AuCl_3 \quad \text{and} \quad Pt + 4[Cl] \rightarrow PtCl_4

10.6 USES OF NITRIC ACID

The uses in this study of compounds chapter pair each application with the property behind it.

Use	Property behind it
Etching designs on copper and brassware	Acts as a solvent for most metals except noble metals
Purifying gold	Impurities like Cu, Ag, Zn, and Pb dissolve in the acid
Rocket fuel oxidant	Acts as an oxidant for the fuel
Preparing fertilisers	Forms calcium nitrate $Ca(NO_3)_2$ , ammonium nitrate $NH_4NO_3$ , nitro chalk $[NH_4NO_3 + CaCO_3]$ , and basic calcium nitrate $CaO \cdot Ca(NO_3)_2$
Making aqua regia	Aqua regia dissolves noble metals

Industrially, nitric acid is used to make explosives such as TNT, synthetic fibres and plastics, nitrates of potassium, ammonium, and silver, and dyes, drugs, and perfumes.

10.7 TESTS FOR NITRIC ACID AND NITRATES

The tests in this study of compounds chapter recognise nitric acid and its salts through the brown colour of $NO_2$ , and through one distinctive ring test.

1. Heating concentrated nitric acid gives brown fumes of nitrogen dioxide:

4HNO_3 \xrightarrow{\Delta} 2H_2O + 4NO_2 + O_2

2. Heating most nitrates gives reddish-brown $NO_2$ fumes. The general pattern is:

\text{Metallic nitrate} \xrightarrow{\Delta} \text{Metallic oxide} + NO_2 + O_2

3. Adding copper to nitric acid or to an acidified nitrate gives dense reddish-brown $NO_2$ :

$Cu + 4HNO_3 \xrightarrow{\Delta} Cu(NO_3)_2 + 2H_2O + 2NO_2$
$Cu + 4NaNO_3 + 4H_2SO_4 \xrightarrow{\Delta} 4NaHSO_4 + Cu(NO_3)_2 + 2H_2O + 2NO_2$

4. Brown ring test. To a solution of a nitrate or nitric acid, add a freshly prepared saturated solution of ferrous sulphate. Pour concentrated sulphuric acid carefully down the side of the test tube so the two liquids form layers, then cool the tube. A brown ring appears at the junction of the two liquids.

The ring is nitroso ferrous sulphate ( $FeSO_4 \cdot NO$ ). Nitric acid oxidises some ferrous sulphate and is itself reduced to nitric oxide:

$6FeSO_4 + 3H_2SO_4 + 2HNO_3 \rightarrow 3Fe_2(SO_4)_3 + 4H_2O + 2NO$
$FeSO_4 + NO \rightarrow FeSO_4 \cdot NO$

Fresh ferrous sulphate is essential. Exposed to air, it is oxidised to ferric sulphate, which will not give the ring. The ring also decomposes if the tube is disturbed, because the heat breaks down the unstable compound.

Nitrate decomposition on heating falls into four families.

Family	Example equation	Products
Alkali metal nitrates (Na, K)	$2NaNO_3 \xrightarrow{\Delta} 2NaNO_2 + O_2$	Nitrite + oxygen only
Most other nitrates (Ca, Zn, Pb, Cu)	$2Pb(NO_3)_2 \xrightarrow{\Delta} 2PbO + 4NO_2 + O_2$	Oxide + $NO_2$ + $O_2$
Silver and mercury nitrates	$2AgNO_3 \xrightarrow{\Delta} 2Ag + 2NO_2 + O_2$	Metal + $NO_2$ + $O_2$
Ammonium nitrate	$NH_4NO_3 \xrightarrow{\Delta} N_2O + 2H_2O$	Nitrous oxide + steam, no residue

Notice the logic of the families. Sodium and potassium nitrates keep the metal in the nitrite. Most other nitrates leave the metal oxide. Silver and mercury go all the way to the free metal. Ammonium nitrate leaves nothing solid behind.

What to practise from Study of Compounds Class 10 ICSE

Two question blocks close this Class 10 ICSE chapter, and both are worth doing in full.

Intext Questions — 11 questions on the definitions (aqua fortis, aqua regia, fixation of nitrogen), the yellow colour of laboratory nitric acid and its removal, and the reason the preparation must stay below $200^\circ C$ .
EXERCISE — 14 questions on nitrate decomposition families, the brown ring test, and balanced equations for the action of nitric acid on metals and non-metals.

For the official syllabus and examination details, see the CISCE website. The rest of the year’s material sits under the Class 10 ICSE section, and the complete Class 10 books collection lists every subject.

Study of Compounds Class 10 ICSE frequently asked questions

Students working through this Class 10 ICSE chapter ask the questions below most often.

Why does dilute nitric acid not give hydrogen with metals?

Dilute nitric acid is a powerful oxidising agent. The nascent oxygen it releases oxidises the hydrogen to water before the gas can escape. The only case in this chapter is very dilute (about 1%) acid reacting with magnesium or manganese at room temperature, which does give hydrogen.

What is the brown ring test in study of compounds class 10 icse?

The brown ring test identifies nitrate ions. Add a freshly prepared saturated solution of ferrous sulphate to the nitrate solution, then pour concentrated sulphuric acid carefully down the side of the tube and cool it. A brown ring of nitroso ferrous sulphate ( $FeSO_4 \cdot NO$ ) forms at the junction of the two liquids.

Which nitrates do not give nitrogen dioxide on heating?

Sodium and potassium nitrates give only the nitrite and oxygen. Ammonium nitrate gives nitrous oxide and steam. Every other nitrate in this chapter — calcium, zinc, lead, copper, silver, and mercury — gives nitrogen dioxide on heating.

Why is laboratory nitric acid yellow and how is the colour removed?

The yellow colour comes from dissolved nitrogen dioxide, formed when part of the acid decomposes in the heat. Bubble dry air or carbon dioxide through the acid to drive out the $NO_2$ , or add excess water so the gas dissolves.

What is aqua regia and why can it dissolve gold?

Aqua regia is one part concentrated nitric acid mixed with three parts concentrated hydrochloric acid. The nitric acid oxidises the HCl to nascent chlorine, and this nascent chlorine attacks gold to form soluble gold(III) chloride, $AuCl_3$ .

Reference: ICSE Class 10 Chemistry textbook material, chapter 10: Study of Compounds Nitric Acid — CISCE-aligned.

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